Calorimeters and Calorimetry Thermal Physics - Lesson 2 Calorimetry Calorimeters and Calorimetry Measuring the Quantity of Heat Calorimeters and Calorimetry Calorimetry is the science associated with determining the changes in energy of a system by measuring the heat exchanged with the surroundings. Now that sounds very textbooky; but in this last part of Lesson 2, we are going to try to make some meaning of this definition of calorimetry. In physics class (and for some, in chemistry class), calorimetry labs are frequently performed in order to determine the heat of reaction or the heat of fusion or the heat of dissolution or even the specific heat capacity of a metal. These types of labs are rather popular because the equipment is relatively inexpensive and the measurements are usually straightforward. In such labs, a calorimeter is used. A calorimeter is a device used to measure the quantity of heat transferred to or from an object. Most students likely do not remember using such a fancy piece of equipment known as a calorimeter. Fear not; the reason for the lack of memory is not a sign of early Alzheimer's.

Rather, it is because the calorimeter used in high school science labs is more commonly referred to as a Styrofoam cup. It is a coffee cup calorimeter - usually filled with water. The more sophisticated cases include a lid on the cup with an inserted thermometer and maybe even a stirrer. Coffee Cup Calorimetry So how can such simple equipment be used to measure the quantity of heat gained or lost by a system? We have learned on the previous page , that water will change its temperature when it gains or loses energy. And in fact, the quantity of energy gained or lost is given by the equation Q = mwater•Cwater•ΔTwater where Cwater is 4.18 J/g/°C. So if the mass of water and the temperature change of the water in the coffee cup calorimeter can be measured, the quantity of energy gained or lost by the water can be calculated. The assumption behind the science of calorimetry is that the energy gained or lost by the water is equal to the energy lost or gained by the object under study.

So if an attempt is being made to determine the specific heat of fusion of ice using a coffee cup calorimeter, then the assumption is that the energy gained by the ice when melting is equal to the energy lost by the surrounding water. It is assumed that there is a heat exchange between the iceand the water in the cup and that no other objects are involved in the heat exchanged. This statement could be placed in equation form as Qice = - Qsurroundings = -Qcalorimeter The role of the Styrofoam in a coffee cup calorimeter is that it reduces the amount of heat exchange between the water in the coffee cup and the surrounding air. The value of a lid on the coffee cup is that it also reduces the amount of heat exchange between the water and the surrounding air. The more that these other heat exchanges are reduced, the more true that the above mathematical equation will be. Any error analysis of a calorimetry experiment must take into consideration the flow of heat from system to calorimeter to other parts of the surroundings. And any design of a calorimeter experiment must give attention to reducing the exchanges of heat between the calorimeter contents and the surroundings.

Bomb Calorimetry The coffee cup calorimeters used in high school science labs provides students with a worthwhile exercise in calorimetry. But at the professional level, a cheap Styrofoam cup and a thermometer isn't going to assist a commercial food manufacturer in determining the Calorie content of their products. For situations in which exactness and accuracy is at stake, a more expensive calorimeter is needed. Chemists often use a device known as a bomb calorimeter to measure the heat exchanges associated with chemical reactions, especially combustion reactions. Having little to nothing to do with bombs of the military variety, a bomb calorimeter includes a reaction chamber where the reaction (usually a combustion reaction) takes place. The reaction chamber is a strong vessel that can withstand the intense pressure of heated gases with exploding. The chamber is typically filled with mostly oxygen gas and the fuel. An electrical circuit is wired into the chamber in order to electrically ignite the contents in order to perform a study of the heat released upon combustion. The reaction chamber is surrounded by a jacket of water with a thermometer inserted. The heat released from the chamber warms the water-filled jacket, allowing a scientist to determine the quantity of energy released by the reaction.

Source: Wikimedia Common s; thanks to Lisdavid89 . Solving Calorimetry Problems Now let's look at a few examples of how a coffee cup calorimeter can be used as a tool to answer some typical lab questions. The next three examples are all based on laboratory experiments involving calorimetry. Example Problem 1: A physics class has been assigned the task of determining an experimental value for the heat of fusion of ice. Anna Litical and Noah Formula dry and mass out 25.8-gram of ice and place it into a coffee cup with 100.0 g of water at 35.4°C. They place a lid on the coffee cup and insert a thermometer. After several minutes, the ice has completely melted and the water temperature has lowered to 18.1°C. What is their experimental value for the specific heat of fusion of ice? The basis for the solution to this problem is the recognition that the quantity of energy lost by the water when cooling is equal to the quantity of energy required to melt the ice.

In equation form, this could be stated as Qice = -Qcalorimeter (The negative sign indicates that the ice is gaining energy and the water in the calorimeter is losing energy.) Here the calorimeter (as in the Qcalorimeter term) is considered to be the water in the coffee cup. Since the mass of this water and its temperature change are known, the value of Qcalorimeter can be determined. Qcalorimeter = m•C•ΔT Qcalorimeter = (100.0 g)•(4.18 J/g/°C)•(18.1°C - 35.4°C) Qcalorimeter = -7231.4 J The negative sign indicates that the water lost energy. The assumption is that this energy lost by the water is equal to the quantity of energy gained by the ice. So Qice = +7231.4 J. (The positive sign indicates an energy gain.) This value can be used with the equation from the previous page to determine the heat of fusion of the ice.

Qice = mice•ΔHfusion-ice +7231.4 J = (25.8 g)•ΔHfusion-ice ΔHfusion-ice = (+7231.4 J)/(25.8 g) ΔHfusion-ice = 2.80x102 J/g (rounded to two significant figures) Example Problem 2: A chemistry student dissolves 4.51 grams of sodium hydroxide in 100.0 mL of water at 19.5°C (in a calorimeter cup). As the sodium hydroxide dissolves, the temperature of the surrounding water increases to 31.7°C. Determine the heat of solution of the sodium hydroxide in J/g. Once more, the solution to this problem is based on the recognition that the quantity of energy released when sodium hydroxide dissolves is equal to the quantity of energy absorbed by the water in the calorimeter.

In equation form, this could be stated as QNaOH dissolving = -Qcalorimeter (The negative sign indicates that the NaOH is losing energy and the water in the calorimeter is gaining energy.) Since the mass and temperature change of the water have been measured, the energy gained by the water (calorimeter) can be determined. Qcalorimeter = m•C•ΔT Qcalorimeter = (100.0 g)•(4.18 J/g/°C)•(31.7°C - 19.5°C) Qcalorimeter = 5099.6 J The assumption is that this energy gained by the water is equal to the quantity of energy released by the sodium hydroxide when dissolving. So QNaOH-dissolving = -5099.6 J. (The negative sign indicates an energy lost.) This quantity is the amount of heat released when dissolving 4.51 grams of the sodium hydroxide. When the heat of solution is determined on a per gram basis, this 5099.6 J of energy must be divided by the mass of sodium hydroxide that is being dissolved.

ΔHsolution = QNaOH-dissolving / mNaOH ΔHsolution = (-5099.6 J) / (4.51 g) ΔHsolution = -1130.7 J/g ΔHsolution = -1.13x103 J/g (rounded to three significant figures) Example Problem 3: A large paraffin candle has a mass of 96.83 gram. A metal cup with 100.0 mL of water at 16.2°C absorbs the heat from the burning candle and increases its temperature to 35.7°C. Once the burning is ceased, the temperature of the water was 35.7°C and the paraffin had a mass of 96.14 gram. Determine the heat of combustion of paraffin in kJ/gram. GIVEN: density of water = 1.0 g/mL. As is always the case, calorimetry is based on the assumption that all the heat lost by the system is gained by the surroundings. It is assumed that the surroundings is the water that undergoes the temperature change.

In equation form, it could be stated that Qparaffin = -Qwater Since the mass and temperature change of the water are known, the energy gained by the water in the calorimeter can be determined. Qcalorimeter = m•C•ΔT Qcalorimeter = (100.0 g)•(4.18 J/g/°C)•(35.7°C - 16.2°C) Qcalorimeter = 8151 J The paraffin released 8151 J or 8.151 kJ of energy when burned. This is based on the burning of 0.69 gram (96.83 g - 96.14 g).

To determine the heat of combustion on a per gram basis, the Qparaffin value (-8.151 kJ) must be divided by the mass of paraffin burned: ΔHcombustion - paraffin = (-8.151 kJ) / (0.69 g) ΔHcombustion - paraffin = -11.813 kJ/g ΔHcombustion - paraffin = -12 kJ/g (rounded to two significant digits) Check Your Understanding 1. Consider Example Problem 3 above. Identify as many sources of error as you can. For each source indicate the direction of error that would have resulted. That is, identify whether the error would have caused the experimentally derived value to be less than or more than the accepted value. Answer: Answers will vary. Three common choices include: A. Energy is transferred from the water to the surrounding air. This would cause the experimental value to be less than the accepted value since this energy is not contributing to the water's temperature change. B. Energy is being absorbed by the metal cup as the metal also encounters a temperature change. This would cause since this energy is not being accounted for in the calculations.

C. Some of the energy released by the burning candle fails to warm either the cup or the water. This energy simply warms the surrounding air. Failure to account for this energy would cause the experimental value to be less than the accepted value. 2. A 2.15-gram cashew nut is burned. The heat released raises the temperature of a 100.0-gram sample of water from 18.2°C to 31.5°C. The mass of the nut after the experiment is 1.78 grams. Determine the calorie content of the nut in Calories/gram. Assume that the water is only able to absorb 25% of the heat released by the burning nut. Given 1.00 Calorie=4.18 kJ. Calorimetry - Measuring Heats of Reactions Calorimetry The picture shows reading the sensitive thermometer while working with the P6310 Bomb Calorimeter . The bomb is inside the tank. Since volume does not change, a bomb calorimeter measures the heat evolved under constant volume, qv, qv = C * dT, where dT is the temperature increase.

The qv so measured is also called the change in internal energy, dE. Note that dE = qv = C * dT Example 2 When a 10.0-watt (J/s) heater is used to heat a bomb calorimeter, its temperature increases by 3.0 K in 5.0 min. Calculate the heat capacity of the calorimeter. formula: Energy used = 10 * 5 * 60 = 3000 J. Thus, C = 3000 / 3.0 = 1000 J/K, This is equivalent to (1000 J/K) / (75.2 J/(K mol)) = 13.3 moles of water. The following related examples illustrate the application of bomb calorimeter for the measurement of dE, and the derivation of dH. Example 3 A calorimeter with heat capacity equivalent to having 13.3 moles of water is used to measure the heat of combustion from 0.303 g of sugar (C12H22O11). The temperature increase was found to be 5.0 K.

Calculate the heat released, the amount of heat released by 1.0 g, and 1.0 mole of sugar. formula: dE = qv = s * m * dT s, specific heat; m mass. Solution: qv = 13.3 mol * 75.2 J/(K mol) * 5.0 K = 5000 J The amount of heat released by 1.0 g would be, 5000 J/0.303 g = 16.5 kJ / g Since the molecular weight of sugar is 342.3 g/mol, the amount of heat released by 1.0 mole would be 16.5 kJ / g * 342.3 g/mol = 5648 kJ/mol. Example 4 The heat released by one mole of sugar from a bomb calorimeter experiment is 5648 kJ/mol. Calculate the enthalpy of combustion per mole of sugar.

Solution The balanced chemical reaction equation is C12H22O11(s) + 12 O2(g) = 12 CO2(g) + 11 H2O(l) Since the total numbers of moles of gases (12) before the reaction is the same as that after the reaction, dn = 0 dH = dE = 5648 kJ/mol Discussion This is the ideal amount of energy released when a mole of sugar is utilized by a living creature such as a person. Example 5 A table of thermodynamic data gives dHf = -285.8 kJ/mol for water. A bomb calorimeter measurement gives the heat of combustion for H2 as -282.0 kJ/mol. Estimate the error of the enthalpy measurement. Solution Reinterpret the problem, we have H2(g) + 0.5 O2(g) = H2O(l), dE = 282.0 kJ/mol.

Furthermore, The error is (285.8-285.7)/285.8 = 0.03% Discussion More heat is giving of if the reaction is carried out at constant pressure, since the P-V work (1.5 R T) due to the compression of 1.5 moles of gases in the reactants would contribute to dH. If 1.0 mole water is decomposed by electrolysis at constant pressure, we must supply an amount of energy equivalent to enthalpy change, dH, a little more than internal energy, dE. More energy must be supplied to perform the P-V work to be done by the products (H2 and O2). The heat capacity of the calorimeter can also be determined by burning an exactly known amount of a standard substance, whose enthalpy of combustion has been determined. Benzoic acid, C7H6O2, is one such standard. The problem below illustrates the calculations. Example 6 When 0.1025 g of benzoic acid was burnt in a bomb calorimeter the temperature of the calorimeter increased by 2.165° C.

For benzoic acid dH°comb = -3227 kJ mol-1. Calculate the heat capacity of the calorimeter.

Solution: The equation for the combustion is, C7H6O2(s) + 7.5 O2(g) -> 7CO2(g) + 3H2O(l), dH° = 3227 kJ Since 7.5 moles of O2 gas is needed, and 7 moles of CO2 is produced, some pressure-volume work is done, to the calorimeter: P V = dn R T, where dn = (7 - 7.5) = - 0.5 mol dE = dH - dn R T = - 3227 - (-0.5*8.314298*298) = - 3226 kJ/mol (a small correction) The amount of heat produced by 0.1025 g benzoic acid is q = 0.1025/122.13 mol x 3226 (kJ/mol) = 2.680 kJ Thus, the heat capacity is C = q / dT = 2.680 / 2.165 = 1.238 kJ / K

.

After the heat capacity is determined, the calorimeter is ready to be used to measure the enthalpy of combustion of other substances. Example 7 When 0.7022 g of oxalic acid (C2O4H2) is burnt in the calorimeter under the same conditions as Example 6, the temperature increased by 1.602°C. The heat capacity of the calorimeter is 1.238 kJ/K. Calculate dH°comb.

Solution: The balanced equation and various quantities calculated are given in a logical order below: C2O4H2(s) + 0.5 O2(g) -> 2 CO2 (g) + H2O(l) dn = 1.5 q = C dT = 1.238*1.602 = 1.984 kJ n of oxalic acid = 0.7022/90 = 0.00780 mol dE = -1.984 / 0.00780 = -354.4 kJ/mol dH = dE + dn R T = -254.4 kJ + 1.5 mol * 0.008314 kJ/(K mol)* 298 K = -250.6 kJ/mol Discussion: Can the standard enthalpy of formation dH°f for oxalic acid be calculated? What additional data are required? Example 8 Calculate the enthalpy of formation of oxalic acid, for which the enthalpy of combustion is -251 kJ/mol.

Solution: The following data were looked up from thermodynamic data , dH°f (CO2(g)) = -393.5 kJ/mol dH°f (H2O(l)) = -285.8 kJ/mol Since C2O4H2(s) + 0.5 O2(g) -> 2 CO2 (g) + H2O(l) dH°comb = 2 dH°f(CO2(g)) + dH°f(H2O(l)) - dH°f(oxalic acid) -250.6 = 2 * (-393.5) + (-285.8) - dH°f(oxalic acid) Therefore dH°f(oxalic acid) = 2 * (-393.5) + (-285.8) - (-250.6) = -822.2 kJ/mol Discussion: It's important to know what additional data is required for problem solving. It's equally important to know where to look for them.

We have given an introduction to calorimetry here, and for a detailed laboratory instruction, consult the CHEM120L manual and the Oxygen Bomb Calorimetry which has several pages giving the details of the measurement. In particular, how to measure the temperature difference is given in Calorimetry Data Analysis . Confidence Building Questions A microwave oven operates at 500 watt for 2 minutes. How much electric energy (J) is consumed? Skill: Calculate amount of electric energy for heat capacity measurement. Review Skills. A cup of (200 g) water heated in a microwave oven had a temperature increase of 60 K. How much energy (in kJ) was absorbed by the water? Recall that heat capacity of water is 4.184 J/(g K). Skill: Calculate the amount of energy when the heat capacity and temperature increase is known, a preparation for enthalpy evaluation. The temperature of a calorimeter increased 0.25 K when 50.0 J of energy was absorbed. What is its heat capacity (kJ/K)? Skill: Measure the heat capacity of a calorimeter.

When 100 mL of 0.50 M NH3 is neutralized by 300 mL of a solution containing excess HCl, the solution temperature increases 1.6 K. The experiment is performed in a styrofoam cup, and there is no energy loss. Calculate the heat of reaction, dH. Discussion: 2678 J was generated by 0.05 mol NH3 (0.100 L * 0.5 mol/L) How much heat (kJ) should be generated if 1.0 mole NH3 is neutralized by dilute HCl solution, if 2675 J was generated when 0.05 mol NH3 (0.100 L * 0.5 mol/L) was neutralized as described in the previous problem? Discussion: In fact, 53.6 kJ/mol is called the molar heat of neutralization of NH3. Combustion of 0.92 g toluene, C7H8 in a bomb calorimeter (Cv equivalent to 1003 g of water) raised its temperature by 9.3 K.

Estimate dH Some details q = 39040 J; dE = 3904 kJ; dH = -3909 kJ/mol. When 0.1234 g of oxalic acid (C2O4H2) is burnt in the calorimeter under the same conditions as Example 6, the temperature increased by 2.314°C. Calculate the heat capacity of the calorimeter. Skill Using a standard to determine the heat capacity of a calorimeter. When 0.2006 g of oxalic acid (C2O4H2) is burnt in the calorimeter whose heat capacity is 1.238 kJ/K. the temperature increased by 0.4577 K. Calculate dH°comb.

Skill: Bomb Calorimetry BOMB CALORIMETRY 1.ï¿½ Purpose of Bomb Calorimetry Experiments Bomb calorimetry is used to determine the enthalpy of combustion, DcombH, for hydrocarbons: CxHYOz (s)ï¿½ +ï¿½ (2X+Y/2-Z)/2 O2 (g)ï¿½ ï¿½ï¿½ X CO2 (g) + Y H2O (l) Since combustion reactions are usually exothermic (give off heat), DcombH is typically negative.ï¿½ (However, be aware that older literature defines the "heat of combustion" as -DcombH, so as to avoid compiling tables of negative numbers!) 2.ï¿½ Construction of a Bomb Calorimeter The bomb calorimeter consist primarily of the sample, oxygen, the stainless steel bomb, and water.

The dewar prevents heat flow from the calorimeter to the rest of the universe, i.e., qcalorimeter = 0 Since the bomb is made from stainless steel, the combustion reaction occurs at constant volume and there is no work, i.e., wcalorimeter = -ï¿½ p dV = 0 Thus, the change in internal energy, DU, for the calorimeter is zero DUcalorimeter = qcalorimeter + wcalorimeter = 0 The thermodynamic interpretation of this equation is that the calorimeter is isolated from the rest of the universe. 3.ï¿½ DU and DH in a Bomb Calorimeter 3.ï¿½ A.ï¿½ Internal energy change DU Since the calorimeter is isolated from the rest of the universe, we can define the reactants (sample and oxygen) to be the system and the rest of the calorimeter (bomb and water) to be the surroundings.

The change in internal energy of the reactants upon combustion can be calculated from Since the process if constant volume, dV=0.ï¿½ Thus, recognizing the definition of heat capacity Cv yields Assuming Cv to be independent of T over small temperature ranges, this expression can be integrated to give where Cv is the heat capacity of the surroundings, i.e., the water and the bomb. 3.ï¿½ B.ï¿½ Enthalphy change DH By definition of enthalpy Since there is very little expansion work done by condensed phases, D(pV) ï¿½ 0 for solids and liquids.ï¿½ Assuming the gas to be ideal yields 3.ï¿½ C.ï¿½ Intuitive difference between DU and DH Recall that DU=qv is the heat flow under constant volume conditions, whereas DH=qpis the heat flow under constant pressure conditions.ï¿½ The difference between these two situations is that pV work can be done under constant pressure conditions, whereas no pV work is done under constant volume conditions. Consider the case where Dngas > 0.

i.e., the system expands during the reaction.ï¿½ The same amount of energy is released by the reaction under both sets of conditions.ï¿½ However, some of the energy is released in the form of work at constant pressure; thus, the heat released will be less than at constant volume.ï¿½ Mathematically, In the case where Dngas < 0.

i.e., the system contracts during the reaction, the surroundings does work on the system.ï¿½ Thus, this work is available for energy release from the system back to the surroundings in the form of heat.ï¿½ Mathematically, These cases can be depicted pictorially as follows: -DH > -DU or DH < DU 4.ï¿½ Calibration of the Calorimeter 4.ï¿½ A.ï¿½ Estimating Cv The heat capacity of the bomb calorimeter can be estimated by considering the calorimeter to be composed of 450 g water and 750 g stainless steel.ï¿½ Knowing the specific heat capacity of water to be 1 cal/gï¿½K and estimating the specific heat capacity of steel to be 0.1 cal/gï¿½K yields 4.ï¿½ B.ï¿½ Measuring Cv For accurate work, the heat capacity of the calorimeter must be measured.ï¿½ This is done by depositing a known amount of energy into the calorimeter and observing the temperature increase.ï¿½ The two most common methods for measuring Cv are 1.ï¿½ Burning a standard with known DU, e.g., benzoic

 acid.

mbenzoic acid DUbenzoic acid = mbenzoic acid -6318 cal/gï¿½K = -Cv DT 2.ï¿½ Doing electrical work by passing current though a resistor. DU = w + q = Vï¿½Iï¿½t + 0 = Cv DT 5.ï¿½ Corrections in Bomb Calorimetry 5.ï¿½ A.ï¿½ Combustion of fuse Nickel and iron fuses can burn according to Ni + The heat released by combustion of the fuse is accounted for by recognizing that DU = DUsampleï¿½msample + DUburned fuseï¿½mburned fuse = -Cv DT where the mass of the burned fuse is determined by weighing the fuse before and after firing the bomb.

5.ï¿½ B.ï¿½ Nonadiabaticity of calorimeter A bomb calorimeter is only approximately adiabatic.ï¿½ In reality, there is a small heat leak through the dewar (qcalorimeter ï¿½ 0) and the stirrer does work on the calorimeter (wcalorimeter ï¿½ 0).ï¿½ Nonadiabaticity is corrected for with an empirical radiative correction, RC.

The time at which the bomb is considered to be fired is the time that makes the areas indicated in the above figure equal.ï¿½ For the Parr calorimeter, this is estimated to be at t = 7 minutes.ï¿½ Thus, the temperature at t = 6 minutes must be extrapolated forward 1 minute by the pre-firing slope, and the temperature at t = 12 minutes must be extrapolated backward 5 minutes by the post-firing slope.ï¿½ Mathematically, this is done as follows 5.ï¿½ C.ï¿½ Nitric acid formation At high temperatures, nitrogen can form nitric acid in the presence of oxygen and water.ï¿½ (This reaction also occurs in automobile engines and is partially responsible for smog production.) N2 + O2 + H2O ï¿½ 2HNO3 Flushing the bomb with oxygen prior to firing, thereby displacing all nitrogen, eliminates nitric acid formation.

6.ï¿½ Application of DcombH In addition to measuring the energy release of one particular reaction, calorimetry is an important tool for determining the enthalpy of formation for the compound under study.ï¿½ This information can then be applied to any reaction involving the compound. The enthalpy of combustion for the reaction DcombH(CxHyOz) = v(CxHyOz)DfHï¿½(CxHyOz) + v(O2)DfHï¿½(O2) + v(CO2)DfHï¿½(CO2) + v(H2O)DfHï¿½(H2O) where v(i) is the stoichiometric coefficient of i.ï¿½ Since DfHï¿½(CxHyOz) and DfHï¿½(H2O) are known (and DfHï¿½(O2) equals zero), measurement of DcombH(CxHyOz) allows calcualtion of DfHï¿½(CxHyOz).

7.ï¿½ Other Types of Calorimeters There are many kinds of calorimeters, each designed for measuring the heat released by a particular chemical process.ï¿½ Some examples include: Flame Calorimeter The combustible gas is metered into the calorimeter.ï¿½ Temperatures of all reactants must be controlled.ï¿½ Since the reaction occurs at constant pressure, DcombH is measured directly. Solution Calorimeter Reactants are initially separated.ï¿½ The temperature change is measured when they are allowed to mix.ï¿½ Quantities that can be determined include DmixH, DdilutionH, and DsolvationH. Calorimeter design is very tricky, especially for processes involving very small energy changes, e.g., protein folding, or energy changes on top of a large background, e.g., excess heat from "cold fusion".ï¿½ Heat leaks must be minimized, and all other heat generating processes must be accounted for. Indirect calorimetry | definition of indirect calorimetry by Medical dictionary Indirect calorimetry | definition of indirect calorimetry by Medical dictionary [kal″o-rim´ĕ-tre] measurement of the heat eliminated or stored in any system.

direct calorimetry measurement of the amount of heat produced by a subject enclosed within a small chamber. indirect calorimetry measurement of the amount of heat produced by a subject by determination of the amount of oxygen consumed and the amount of carbon dioxide eliminated. in·di·rect cal·o·rim·e·try determination of heat production of an oxidation reaction by measuring uptake of oxygen and/or liberation of carbon dioxide and nitrogen excretion and then calculating the amount of heat produced. indirect calorimetry the measurement of the amount of heat generated in an oxidation reaction by determining the intake or consumption of oxygen or by measuring the amount of carbon dioxide or nitrogen released and translating these quantities into a heat equivalent. Compare direct calorimetry . in·di·rect ca·lor·i·me·try (in'di-rekt' kal'ŏr-im'ĕ-trē) Determination of heat production of an oxidation reaction by measuring uptake of oxygen or liberation of carbon dioxide and nitrogen excretion. calorimetry What is a Calorimeter? (with picture) What is a Calorimeter?

Last Modified Date: 17 January 2017 Copyright Protected: Top 10 facts about the world As one of the devices used in the science of calorimetry , the calorimeter plays an important role in measuring heat capacity in a number of secure environments, as well as the physical changes that occur with the incidence of specific chemical reactions . The word is derived from the Latin term calor, which literally translates to "heat." While most people understand the concept of calories as they relate to weight gain or loss, fewer individuals are aware that calorie content is determined by the ability of the substance to generate heat or energy as a result of the consumption. What the calorimeter actually does is identify the rate and scale of the production of that heat or energy and translate it into numerical designations that can be used to adequately measure the production and the assimilation of heat and energy. When thinking in terms of counting calories, therefore, what is actually happening is thinking in terms of how much heat and energy will be generated and expended as a result. Ad Calorie measurement has several important applications. One of the most common has to do with health matters.

Because the calorie content of some foods is judged to be quite high, the result is that the energy or heat created by the consumption of the foods must make best use of that produced energy. If there is no efficient use of the calories, then the result is energy stored in the form of fat for later consumption. That is why the calorimeter is so important. Proper measuring of calories allows people to more properly plan their food intake to ensure that the body has enough heat or energy to function, but does not store excess energy as fat. There are a number of different types of the measuring tools that help to measure energy and heat production, including differential scanning, accelerated rate, and isothermal calorimeters. Each of these are used in some way to measure calories within different environments. They may be used to note the change in temperature in a substance when a liquid is introduced, for example. Other factors, such as a change in weather conditions within the controlled environment, also yield data about the change in heat and energy production. All these types of research and scientific experiment help to refine the way the concept of calories are used within the scientific community as well as in industry and in daily life.

While the calorimeter is not a device that is found in the home, the ongoing investigation into the way specific heat capacities are determined does make an impact on the way people live, work, and eat. Use of these precise instruments help to make the science of energy and heat production and usage more efficient, as well as make the use of calories more productive for society in general. Ad Coffee Cup Calorimetry and Bomb Calorimetry Coffee Cup Calorimetry and Bomb Calorimetry Coffee Cup Calorimetry and Bomb Calorimetry Measurement of Heat Flow & Enthalpy Change This is a cross section of a bomb calorimeter, one of the most common types of calorimeters used to measure heat flow in a reaction. Encyclopaedia Britannica/UIG, Getty Images A calorimeter is a device used to measure the quantity of heat flow in a chemical reaction. Two of the most common types of calorimeters are the coffee cup calorimeter and the bomb calorimeter. Coffee Cup Calorimeter A coffee cup calorimeter is essentially a polystyrene (Styrofoam) cup with a lid.

The cup is partially filled with a known volume of water and a thermometer is inserted through the lid of the cup so that its bulb is below the water surface. When a chemical reaction occurs in the coffee cup calorimeter, the heat of the reaction if absorbed by the water. The change in the water temperature is used to calculate the amount of heat that has been absorbed (used to make products, so water temperature decreases) or evolved (lost to the water, so its temperature increases) in the reaction. Heat flow is calculated using the relation: q = (specific heat) x m x Δt where q is heat flow, m is mass in grams , and Δt is the change in temperature. continue reading below our video 10 Best Universities in the United States The specific heat is the amount of heat required to raise the temperature of 1 gram of a substance 1 degree Celsius. The specific heat of water is 4.18 J/(g·°C). For example, consider a chemical reaction which occurs in 200 grams of water with an initial temperature of 25.0°C. The reaction is allowed to proceed in the coffee cup calorimeter.

As a result of the reaction, the temperature of the water changes to 31.0°C. The heat flow is calculated: qwater = 4.18 J/(g·°C) x 200 g x (31.0°C - 25.0°C) qwater = +5.0 x 103 J In other words, the products of the reaction evolved 5000 J of heat, which was lost to the water. The enthalpy change , ΔH, for the reaction is equal in magnitude but opposite in sign to the heat flow for the water: ΔHreaction = -(qwater) Recall that for an exothermic reaction, ΔH < 0; qwater is positive. The water absorbs heat from the reaction and an increase in temperature is seen. For an endothermic reaction, ΔH > 0; qwater is negative. The water supplies heat for the reaction and a decrease in temperature is seen. Bomb Calorimeter A coffee cup calorimeter is great for measuring heat flow in a solution, but it can't be used for reactions which involve gases, since they would escape from the cup.

The coffee cup calorimeter can't be used for high temperature reactions, either, since these would melt the cup. A bomb calorimeter is used to measure heat flows for gases and high temperature reactions. A bomb calorimeter works in the same manner as a coffee cup calorimeter, with one big difference. In a coffee cup calorimeter, the reaction takes place in the water. In a bomb calorimeter , the reaction takes place in a sealed metal container, which is placed in the water in an insulated container. Heat flow from the reaction crosses the walls of the sealed container to the water. The temperature difference of the water is measured, just as it was for a coffee cup calorimeter. Analysis of the heat flow is a bit more complex than it was for the coffee cup calorimeter because the heat flow into the metal parts of the calorimeter must be taken into account: qreaction = - (qwater + qbomb) where qwater = 4.18 J/(g·°C) x mwater x Δt The bomb has a fixed mass and specific heat. The mass of the bomb multiplied by its specific heat is sometimes termed the calorimeter constant, denoted by the symbol C with units of joules per degree Celsius.

The calorimeter constant is determined experimentally and will vary from one calorimeter to the next. The heat flow of the bomb is: qbomb = C x Δt Once the calorimeter constant is known, calculating heat flow is a simple matter. The pressure within a bomb calorimeter often changes during a reaction, so the heat flow may not be equal in magnitude to the enthalpy change. Calorimetry Calorimetry Calorimetry is the measurement of the quantity of heat exchanged. For example, if the energy from an exothermic chemical reaction is absorbed in a container of water, the change in temperature of the water provides a measure of the amount of heat added. Calorimeters are used to determine the energy content of foods by burning the foods in an oxygen atmosphere and measuring the energy yield in terms of the increase in temperature of the calorimeter. Calorimeters can also be used to measure the specific heat of a substance. The materials involved in the calorimetry are modeled here as a volume of water, a source of heat which is characterized by its equivalent water mass, and the container or calorimeter with its mass and specific heat.

Thermal equilibrium is assumed after the experiment so that the temperature change can be used to calculate the energy released. Mass of water = What does a calorimeter directly measure? chang... - OpenStudy What does a calorimeter directly measure? change in temperature kinetic energy specific heat radiation Physics

Question: What does a Calorimeter measure?
Answer: Heat